I’m melting!!
THERMODYNAMICS
Chemistry 332
Name _______________________________ Class _______ Date _______________
A Heating Curve
1.
The heating curve shown in the figure is a plot of temperature vs. time.
It represents the heating of what is initially ice at -10oC at a constant rate
of heat transfer.
1. a. What phase or phases are
present during segment A?
1 a.
b. What is happening to the energy being
absorbed from the heat source?
(Answer in terms of potential
and/or kinetic energy.)
b.
c. What phase change, if any, is taking
place?
c.
2. a. What phase or phases are present
during segment B?
2 a.
b. What is happening to the energy being
absorbed?
b.
c. What phase change, if any, is taking
place?
c.
d. What is the significance of the
temperature 00 C ?
d.
1
A Heating Curve (Continued)
3. a. What phase or phases are present
during segment C?
b. What is happening to the energy being
absorbed?
b.
c. What phase change, if any, is taking
place?
c.
4. a. What phase or phases are present
during segment D?
5.
3 a.
4 a.
b. What is happening to the energy being
absorbed?
b.
c. What phase change, if any, is taking
place?
c.
d. What is the significance of the
temperature 10 00 C ?
d.
What would you expect to happen if
the heating were continued?
5.
2
Name
Specific Heats
Water
Aluminum
Sand
Copper
4.18 J/g c
9.20 J /g c
6.64 J/g c
3.80 J/g c
1.0 cal/g c
2.2 cal/g c
1.6 cal/g c
.91 cal/g c
1. The temperature of 0.5 kg of water is raised from 25 degrees Celsius to
85 degrees Celsius. How many joules of heat energy were required to do
this?
2. If a 2 kg bag of sand is left in the sun all day, its temperature goes from
12 degrees to 47 degrees Celsius. How much heat energy did the sand
absorb?
3. A 1.5 kg roll of aluminum foil cools from 100oC to 72o. How much heat
is given off?
4. A 20 kg ingot of copper is cooled from 200o to 100o by dropping it into
water. How much heat will the copper give off?
5. How many joules of energy will the water gain?
3
Heat Activity
Part A
Place 100g of sand into a 150 ml beaker. Place 100g of water into another 150 ml
beaker. Place 100g of acetone into a third 150 ml beaker. Place all three beakers on your
hot plate which is at setting #3. Heat and record the temperatures for 10 minutes.
Remove all three beakers from the hot plate after 10 minutes. Let the beakers cool and
record the temp for 10 minutes. Graph your results. Leave the sand and acetone at your
table. Dump the water. Graph your results.
 100g of water = 100ml of water
 100g of acetone = 127 ml of acetone
Part B
Place 100 ml of water into a 400ml beaker. Place 300 ml of water into another
beaker of similar size. Put both beakers onto your hot plate. Keep the hot plate setting
#5. Heat both beakers for 10 minutes and record the temperature changes. Remove
both beakers from the hot plate after 10 minutes. Continue to record the temperatures
for another 10 minutes. Graph your results.
Answer the following questions while you are at the lab:
1. What is the difference between temperature and heat?
2. What units are used to measure temperature? Heat?
3. What devices are used to measure temperature? Heat?
4
PART A
Time (min)
Sand
Acetone
Water
.5
1
1.5
2
2.5
3
3.5
4
4.5
5
5.5
6
6.5
7
7.5
8
8.59
9
9.5
10
After Heating
PART A
Time (min)
10.5
11
11.5
12
12.5
13
13.5
14
14.5
15
15.5
16
16.5
17
17.5
18
18.5
19
19.5
20
Sand
Acetone
Water
5
PART B
Time (min)
Large Beaker
Small Beaker
.5
1
1.5
2
2.5
3
3.5
4
4.5
5
5.5
6
6.5
7
7.5
8
8.5
9
9.5
10
After Heating
PART B
Time (min)
Large Beaker
Small Beaker
10.5
11
11.5
12
12.5
13
13.5
14
14.5
15
15.5
16
16.5
17
17.5
18
18.5
19
19.5
20
6
Heat Activity Questions
Part A
1. List all of the variables that were held constant in part A.
2. Which variable was not held constant?
3. Which substance heated up the quickest when placed on the hot plate?
4. Which substance cooled off the quickest?
Part B
1. List all of the variables that were held constant in part B.
2. Which variable was not held constant?
3. Which beaker heated up the quickest when placed on the hot plate?
4. Which beaker cooled off the quickest?
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332 CHEMISTRY
Name _____________________
PLW – SPECIFIC HEAT LAB
Equations needed:
∆ H =m C ( t2 – t1 )
∆ Hwater = - ( ∆ Hprocess )
1. A 25.0 gram metal sphere at 95.0oC is dropped into an insulated
container with 100.0 ml of water at 24.0oC. The final temperature
of the mixture was 36.1oC. What is the specific heat of the metal?
2. A 125 gram metal object at 100.0oC is dropped into an insulated
container partially filled with 125 ml of water at 24.6oC. The
temperature of the water rises to 38.2oC and then slowly starts to drop.
What is the specific heat of the metal?
3. Why do you think that the temperature of the mixture in question #2
starts to drop after reaching a maximum temperature?
4. Why must the metal be completely covered by cool water in the specific
heat determination?
8
Name ______________________
Lab-Specific Heat
Introduction: We have now seen that heat is a measure of the total energy of
a system. It is dependent on the temperature change of the
system, the mass of the system, and the identity of the system.
Quantitatively, we have see that:
(1)
q = mc∆T
or
heat = mass * specific heat capacity * temperature
We have also seen that since heat cannot be measured directly, we can only measure it by
comparing the heat change of one system to the measurable heat change of another system.
Quantitatively, that means that:
(2)
qlost = qgained
or
heat lost = heat gained
But since heat is directional (into or out of a system) we must include a negative sign to
account for the fact that one system is losing heat while another system gains heat:
(3) qlost = - qgained
or
heat lost = - heat gained
Substituting our above equation (1) for heat in our equation (3) we get:
(4) m1c1∆T1 = - m2c2∆T2
Thus, if we know both masses (m1, m2) and we can measure both changes in temperature
(∆T1, ∆T2) and we know the specific heat of one of the substances (c1), then we can calculate
the specific heat of the other substance (c2). Keep in mind that ∆T is always final temperature
(Tf) – the initial temperature (Ti).
(5) m1c1 ( Tf - Ti ) 1 = - m2c2 (Tf - Ti )2
Procedure:
1) In a graduated cylinder, measure 100 mL (which is 100 grams) of tap water.
2) Pour the water into the Styrofoam cup and take its temperature.
3) The metal is located in the fume hood. The mass of the metal is 200g.
4) Use the thermometer provided to determine the temperature of the hot metal.
5) Using tongs, gently and safely, move the hot metal piece to the water in the
Styrofoam cup. Use the other Styrofoam cup to cover the top. This will reduce
unnecessary heat loss.
6) Stir gently using the thermometer until the temperature stops changing.
7) Record the final temperature of the system.
8) Pour the water down the drain, dry off the metal piece and perform the
calculations.
9
Name _______________________
Data Tables
Water
Mass (g)
Cwater (cal/goC)
Ti water (oC)
Tf water (oC)
Metal
Mass (g)
Cmetal (cal/goC)
Ti metal (oC)
Tf metal (oC)
X
Calculations
A. Determine the heat gained/lost by the water.
B. Determine the heat gained/lost by the metal.
C. Determine the specific heat(heat capacity) of the metal.
D. Determine the % Error.
10
Analysis:
Using the formula developed above (5), and making water #1 and the
metal #2, calculate the specific heat of the unknown metal by solving
for c2 or cmetal.
mwatercwater(Tf - Ti)water = - mmetalcmetal(Tf - Ti)metal
Compare your answer to the table below to see which metal yours is:
Metal
Aluminum
Brass
Copper
Iron
Lead
Tin
Specific Heat
Capacity (cal/goC)
0.215
0.0920
0.0923
0.107
0.0305
0.0545
Conclusion:
Which metal do you have? Is it conclusive? What other property could you
check if you are not sure? What are sources of error in the lab?
11
332 Chemistry
EXPERIMENT – HEAT OF FUSION
A calorimeter is a device used to measure heat. You will use a calorimeter to measure the heat
of fusion per gram of ice and the molar heat of fusion of ice. A calorimeter is simply a container
which is insulated to prevent heat from entering or leaving. A Styrofoam thermal cup makes a
good calorimeter.
Use a graduated cylinder to measure exactly 100 ml of room temperature water. Pour it into a
Styrofoam cup. Record the temperature of the water to the nearest 0.1oC. An ice water bath is
located at the front of the room. Notice the temperature of this ice water bath is 0ºC. Therefore
the ice and water are both at 0ºC. Remove one or two ice cubes from the ice bath and place
them into your cup of water. Be sure to shake any excess liquid off of the cubes before placing
it into your 100 ml of water.
Gently stir the ice and water mixture with the thermometer. Observe the temperature as it goes
down. Do not let the ice cubes melt completely. Remove the remaining ice and shake as much
water as possible back into the cup. Measure and record the volume of water remaining in the
cup. Subtract 100 ml from this volume to get the volume of ice that melted. Since the mass of 1
ml of liquid water is 1 gram you can find the mass of the ice that melted.
Data Table
Initial Volume of Water
__________ mL
Final Volume of Water
__________ mL
Volume of ice that melted
__________ mL
Mass of ice that melted ( 1 mL = 1 g )
__________ grams
Initial Temperature of Water
__________ oC
Final Temperature of Water
__________ oC
Change in Temperature ∆T = Tf - Ti
__________ oC
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Post Lab Questions
1. Calculate the amount of thermal heat energy lost by the water in calories.
q = mc∆T
2. What is the amount of thermal heat energy gained by the ice?
qwater = -qice
3. Are your results similar to those of other lab groups? Why or why not?
4. Use the amount of heat gained by the ice to calculate the Heat of Fusion per gram of water.
To do this, divide the amount of heat gained by the ice by the mass of ice that melted.
5. From your calculations, what should be the units on Heat of Fusion?
6. Are your results similar to those of other lab groups? Why or why not?
7. The accepted value for the Heat of Fusion for water is 80 calories per gram. What is your
percent error?
13
Table of Physical Energy Constants
Heat
Capacity
Calories
(gram)(oC)
Heat of
Fusion
Calories
gram
Heat of
Vaporization
Calories
gram
Melting
Point
o
C
Boiling
Point
o
C
Aluminum
Bromine
Copper
Solid
.22
.0703
.092
Liquid
.25
.11
.10
Gas
.19
.0538
.078
96
15.8
49.0
2520
44.8
1150
660
7
1083
2450
59
2582
Gold
Hydrogen
Iodine
.0306
.57
.0518
.033
.231
.076
.0252
3.42
.0348
15.4
14
14.7
415
107
41
1063
-259
114
2660
-253
184
Magnesium
Mercury
Oxygen
.235
.0334
.336
.288
.0332
.394
.204
.0250
.219
88.8
2.78
3.3
1340
69.2
50.9
650
-39
-219
1117
357
-183
Potassium
Sodium
Water
.178
.295
.50
.18
.333
1
.127
.216
.40
15
27
80.
483
1050
540.
63
98
0
757
889
100
Problems
1. How many calories are gained by 29 grams of water when the temperature of
the water changes from 31oC to 49oC?
2. How much heat will 1.12 grams of water lose if its temperature changes from
31.2oC to 11.1oC?
3. How many calories are required to change 4.0 grams of aluminum from 21 oC
to 38oC?
4. How much heat is required to raise the temperature of 67.2 grams of copper
from 14oC to 92oC?
5. How much heat is required to increase the temperature of 3.50 grams of solid
aluminum 56.0oC?
6. How much heat must 564 grams of water lose if its temperature changes from
92.1oC to 41.0oC?
7. How much heat is required to melt 2.6 grams of copper at 1083oC?
8. How much heat is required to melt 3.16 moles of mercury at -39oC?
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9. How much heat is required to boil 45.8 grams of water at 100oC?
10. How much heat is released when 25 grams of oxygen liquefy at -183oC?
11. How much heat is released as 137 grams of potassium cool from 256 oC to
63oC?
12. A sample of bromine weighting 41.5 grams evaporates at 59 oC. How much
heat is absorbed?
13. 60.4 grams of gold is melted at 1063oC. How much heat is required?
14. If 32.0 grams of steam is heated from 120oC to 242oC, how much heat is
absorbed?
15. How much heat is released as 75 grams of water freeze at 0oC?
16. .48 moles of copper is heated from 1275oC to 1400oC. How much heat is
absorbed?
17. Sodium is sometimes used in the cooling coils of a nuclear reactor because it
stays a liquid over a wide range of temperature. If 319 grams of sodium
has its temperature raised from 99oC to 750oC, how much heat does it
absorb?
18. How much heat would 255 grams of sodium absorb as it boils at 889oC?
19. How much heat would 175 grams of sodium absorb as it melts at 98 oC?
20. How much heat is released as 35.9 grams of magnesium solidify at 650oC?
21. How much heat is released as 42.85 grams of aluminum cools from 3800 oC
to 2500oC?
22. When 117 grams of oxygen is warmed from -210oC to -185oC, how much
heat is absorbed?
23. How much heat is absorbed when 844 grams of water boil at 100oC?
24. How much heat is released when 16.8 grams of mercury freezes at
-39oC?
15
25. How much heat is absorbed as 12 grams of hydrogen warms from 25 oC to
78oC?
26. When 248 grams of a metal at 82oC is placed in 351 grams of water at 15oC a
temperature for both metal and water of 27oC is reached. What is the heat
capacity of the metal?
27. A rock weighing 422 grams is heated to a temperature of 100oC. It is then
placed in 85.0 grams of water at 12oC. The final temperature of the
mixture is 38oC. What is the heat capacity of the rock?
28. How much heat is required to change 78 grams of ice at 0 oc to steam at
165oC?
29. How much heat is released as 55 grams of steam at 132oC cools to water at
48oC?
30. How much heat is required to convert 20.0 grams of mercury at -100oC to
75oC?
31. How much heat is released as 140 grams of gaseous bromine at 150oC is
converted to solid bromine at -60oC?
32. How much heat is absorbed when 55.0 grams of solid aluminum at 450oC is
converted to gas at 2.450oC?
33. How much heat is required to convert 23.0 grams of solid potassium at 63 oC
to 250oC?
34. How much heat is released as 90.0 grams of hydrogen cools from
-100oC to -270oC?
35. How much heat is released as 100 grams of iodine gas at 184oC is cooled to
liquid at 114oC?
36. How much heat is required to convert 67.0 grams of solid potassium at
-20oC to gaseous potassium at 950oC?
37. How much heat is required to convert 20.0 grams of solid sodium at 98 oC to
a gas at 889oC?
16
Name ______________________
Heat Problems
q = cm∆T
cwater = 1 cal/goC
1 cal = energy required to raise 1 g of water by 1oC
1. How much energy (in calories) does it take to heat 50 grams of water from 21oC to 50oC? (
C
water = 1 cal/goC)
2. How much energy (in calories) is released when 2000 grams of water cools from 50oC to
10oC? (Cwater = 1 cal/goC)
3. Lake Michigan contains about 1180 cubic miles of water. On average over the past 25 years,
the temperature difference of the water varies from about 2oC in the winter to about
19oC in the winter. If the total mass of water in Lake Michigan is 5 x 1018 grams, then
how many calories are needed to heat up Lake Michigan? (Cwater = 1 cal/goC)
4. How much heat does it take to raise the temperature of a 250 gram bar of iron
(Ciron = 0.108 cal/goC) from 20o to its melting point of 1535oC?
5. What is the final temperature when 2500 calories is applied to a 250 gram block of aluminum
(CAl = 0.215 cal/goC) at room temperature (22oC)?
6. How much water could be heated from 22oC to 40oC if 5000 calories are applied?
(Cwater = 1 cal/goC)
7. What is the temperature change if 100 calories is applied to a 50 gram block of copper? (ccu =
0.092 cal/ goC)
8. What is the specific heat of a substance if 500 calories makes the temperature of 25 grams of
the substance change from 20oC to 37oC
17
Name ____________________
Heat Transfer and the Thermos
Check out http://www.howstuffworks.com for most of these answers.
1. What are the three different types of heat transfer?
A. _________________________________
Example:
B. _________________________________
Example:
C. _________________________________
Example:
2. A good thermos has to prevent heat transfer of which of these types?
3. Why can’t a thermos ever be “perfect”? Where is most of the heat lost?
4. Cathy and Tom each have identical amounts of identical hot soup at the same temperature.
They each have identical thermoses. Each pours the soup into the thermos identically and
they store the thermoses in identical environments. After a period of time, the temperatures
are measured and Cathy’s soup I hotter than Tom’s soup. When asked to explain this
phenomenon, Cathy says she prepped her thermos before adding the soup by adding boiling
water to her thermos and then dumping it out. How could this simple action keep her soup
hotter longer?
18
5. A diagram of a thermos is shown below. Fill in the boxes to describe what the corresponding
arrow is pointing at.
What is in this area?
What types of heat
transfer does this
cut down on?
What is this part
of this container
usually made
of? (be specific)
What types of
heat transfer
does this cut
down on?
6. How does a thermos “know” to keep hot stuff hot and cold stuff cold? What does a thermos
actually do (or not do, more to the point)?
19
20
SOLID
LIQUID
GAS
Place the letter of each description in the most appropriate region of the Venn diagram above.
Also, label the phase changes marked by arrows above.
a. Take the shape of the container
b. Fill up the volume of the container
c. Have a definite mass
d. Are made up of moving particles
e. Have the highest kinetic energy
f. Have a definite volume at a given temperature
g. Have particles that vibrate and/or spin in place
h. Have the lowest kinetic energy
i. They flow (are called fluids)
j. May be elements, compounds, or mixtures
k. Are most compressible
l. Have a crystal lattice
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THERMODYNAMICS NOTE PAGE
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